1. Chapter 4 Arrangement of Electrons in Atoms

2. The new atomic model Rutherford’s model of the atom was an improvement, but it was incomplete. It did not explain where the atom’s negatively charged electrons are located in the space surrounding it’s positively charged nucleus. In the early 20 th century there was an intimate relationship discovered between light and an atom’s electrons. This discovery led to a new view of the nature of energy, matter, and atomic structure.

3. Properties of Light Before 1900, scientists thought light behaved solely as a wave. Later, they discovered particle-like characteristics in waves. Now how do we describe light? As waves of energy or waves of particles?

4. The Wave Description of Light Visible light is a kind of electromagnetic radiation. Electromagnetic radiation- a form of energy that exhibits wavelike behavior as it travels through space. Electromagnetic Spectrum- all the forms of electromagnetic radiation.

5. All forms of electromagnetic radiation travel at a constant speed of about 3.0 x 10^8 m/s through a vacuum, and at slightly slower speeds through matter. Since air is mostly space the value of 3.0 x 10^8 m/s is also light’s approximate speed through air as well.

6. The significant feature of wave motion is its repetitive nature which can be characterized by the measureable properties of wavelength and frequency. Wavelength- (upside down y) the distance between adjacent waves. Frequency- the # of waves that pass a given point in a specific time (usually one second) One wave per second is called a hertz (Hz) P 92 figures 4-1 and 4-2

7. Frequency and wavelength are mathematically related to each other. c= (^) v C = the speed of light, ^= wavelength, v= frequency. Because c is the same for all electromagnetic radiation, the product of ^v is a constant. So… we know that ^ is inversely proportional to v as the ^ increases, then the v decreases and vise versa.

8. Photoelectric Effect Photoelectric effect- refer to the emission of electrons from a metal when light shines on the metal. the mystery of the photoelectric effect involved the frequency of light striking the metal. For a given metal, no electrons were emitted if the lights frequency was below a certain minimum- no matter how long the light was shone. The wave theory predicted that light of any frequency could supply enough energy to eject an electron.

9. The particle description of light. Quantum- the minimum quantity or energy that can be lost or gained by an atom. E = hv E = energy of a quantum of radiation (in joules) v = frequency of radiation emitted h = 6.626 x 10^-34 J x s (Planck’s constant)

10. Photon- a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. E (photon) = hv Ground state- the lowest energy state of an atom. Excited state- an atom has more potential energy than in its ground state

11. Whenever an excited hydrogen atom falls back from an excited state to its ground state or to a lower energy excited state… it emits a photon of radiation. The energy of this photon is equal to the difference between the atoms initial state and its final state. E(photon)= E 2 – E 1 = hv

12. Bohr Model Orbits – allowed paths for electrons to circle the nucleus. Electrons orbit the nucleus in energy levels. Energy levels can be thought of as rungs on a ladder. You can be on the first or second rung but you can’t stand in between them. The higher rung you are on, the more potential energy you have.

13. Chapter 4 Section 3 ► Electron configuration

14. ► The quantum model of the atom improves on the bohr model because it describes the arrangements of electrons in atoms other than hydrogen. ► Electron configuration- the arrangement of electrons in an atom ► Electrons tend to assume arrangements that have the lowest possible energies. ► The lowest-energy arrangement of electrons is called the elements ground state electron configuration

15. Orbitals ► s orbital = 2 electrons max ► p orbital = 6 electrons max ► d orbital = 10 electrons max ► f orbital = 14 electrons max

16. Rules for electron configuration ► To build up electron configurations for the ground state of any particular atom, first the energy levels of the orbitals are determined. ► Then electrons are added one by one according to three basic rules.

17. The three rules ► the first rule shows the order in which electrons occupy orbitals. ► Aufbau principle- an electron occupies the lowest-energy orbital that can receive it. ► Just like in nature, every thing flows towards the least energy.

18. Atomic orbitals in order of increasing energy ► 1s ► 2s, 2p ► 3s, 3p, 4s, 3d ► 4p, 5s, 4d, 5p, 6s, 4f ► Electrons must first fill the 1s orbital before starting on 2s ► You must fill the 2s orbital before starting the 2p ► Not all orbitals are filled in the order you would think.

19. Rule two ► Pauli exclusion principle- no two electrons in the same atom can have the same set of four quantum numbers. ► Pauli exclusion principle- two electrons in the same orbital must have opposite spins.

20. Rule three ► Hund’s rule- orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all elctrons in singly occupied orbitals must have the same spin. ► Hund’s rule- no orbital in an energy level gets a second electron until all orbitals in an energy level have one, AND all obitals that only have one electron must be spinning them in the same direction.

21. Representing Electron Configuration ► Three methods or notations are used to indicate electron configuration. ► Orbital notation see page 106-107 in book.

22. Elements of the second period ► Acording to the aufbau principle, after H and He wich have only one energy level, the period 2 elements must first fill up the s orbital in the n2 level. ► Highest occupied level- the electron containing energy level with the highest principle quantum number. (outer shell) ► Inner-shell electrons- electrons which are not in the highest occupied energy level.

23. Third Period Elements ► After the outer octet is filled in neon, the next electron enter the s sublevel in the n = 3 ► thus the electron configuration of Na will be 1s^2 2s^2 2p^6 3s^1 ► Rather than having to write all that down… we can use noble gas notation ► Noble gas notation is using the last noble gas on the periodic table as a short-hand then configuring the outer-shell electrons regularly. ► So Na can be written as [Ne]3s^1

24. ► When writing electron configuration, it is ok to write them in numerical order instead of the order of increasing energy. ► For example: calcium could be written as follows ► 1s^2 2s^2 2p^6 3s^2 3p^6 3d 4s^2 ► Scandium (Sc) would be as follows: ► 1s^2 2s^2 2p^6 3s^2 3p^6 3d^1 4s^2 ► OR it’s noble gas configuration [Ar]3d^1 4s^2 ► Electrons of an atom can be counted by adding the exponents.

25. Quiz ► Write the electron configuration for the following elements: ► Lithium (Li) ► Boron (B) ► Carbon (C) ► Write the FULL electron configuration for the following elements: ► Sodium (Na) ► Aluminum (Al) ► Silicon (Si) ► Write the noble gas configuration for #’s 4,5,&6